Bonding and structure

These are the exact learning outcomes Cambridge sets for this topic. The candidate is expected to be able to do each of these on the relevant paper.

1. describe, in simple terms, the lattice structure of a crystalline solid which is: (a) giant ionic, including sodium chloride and magnesium oxide (b) simple molecular, including iodine, buckminsterfullerene C60 and ice (c) giant molecular, including silicon(IV) oxide, graphite and diamond (d) giant metallic, including copper
2. describe, interpret and predict the effect of different types of structure and bonding on the physical properties of substances, including melting point, boiling point, electrical conductivity and solubility
3. deduce the type of structure and bonding present in a substance from given information
4. understand that chemical reactions are accompanied by enthalpy changes and these changes can be exothermic (ΔH is negative) or endothermic (ΔH is positive)
5. construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy
6. define and use the terms: (a) standard conditions (this syllabus assumes that these are 298 K and 101 kPa) shown by ⦵. (b) enthalpy change with particular reference to: reaction, ΔHr , formation, ΔHf , combustion, ΔHc , neutralisation, ΔHneut
7. understand that energy transfers occur during chemical reactions because of the breaking and making of chemical bonds
8. use bond energies (ΔH positive, i.e. bond breaking) to calculate enthalpy change of reaction, ΔHr
9. understand that some bond energies are exact and some bond energies are averages
10. calculate enthalpy changes from appropriate experimental results, including the use of the relationships q = mcΔT and ΔH = –mcΔT/n 5.2 Hess’s law Learning outcomes
11. apply Hess’s law to construct simple energy cycles
12. carry out calculations using cycles and relevant energy terms, including: (a) determining enthalpy changes that cannot be found by direct experiment (b) use of bond energy data
13. calculate oxidation numbers of elements in compounds and ions
14. use changes in oxidation numbers to help balance chemical equations
15. explain and use the terms redox, oxidation, reduction and disproportionation in terms of electron transfer and changes in oxidation number
16. explain and use the terms oxidising agent and reducing agent
17. use a Roman numeral to indicate the magnitude of the oxidation number of an element
18. define Le Chatelier’s principle as: if a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change
19. use Le Chatelier’s principle to deduce qualitatively (from appropriate information) the effects of changes in temperature, concentration, pressure or presence of a catalyst on a system at equilibrium
20. deduce expressions for equilibrium constants in terms of concentrations, Kc
21. use the terms mole fraction and partial pressure
22. deduce expressions for equilibrium constants in terms of partial pressures, Kp (use of the relationship between Kp and Kc is not required)
23. use the Kc and Kp expressions to carry out calculations (such calculations will not require the solving of quadratic equations)
24. calculate the quantities present at equilibrium, given appropriate data
25. state whether changes in temperature, concentration or pressure or the presence of a catalyst affect the value of the equilibrium constant for a reaction
26. describe and explain the conditions used in the Haber process and the Contact process, as examples of the importance of an understanding of dynamic equilibrium in the chemical industry and the application of Le Chatelier’s principle 7.2 Brønsted–Lowry theory of acids and bases Learning outcomes
27. state the names and formulas of the common acids, limited to hydrochloric acid, HCl, sulfuric acid, H2SO4, nitric acid, HNO3, ethanoic acid, CH3COOH
28. state the names and formulas of the common alkalis, limited to sodium hydroxide, NaOH, potassium hydroxide, KOH, ammonia, NH3
29. describe the Brønsted–Lowry theory of acids and bases
30. describe strong acids and strong bases as fully dissociated in aqueous solution and weak acids and weak bases as partially dissociated in aqueous solution
31. appreciate that water has pH of 7, acid solutions pH of below 7 and alkaline solutions pH of above 7
32. explain qualitatively the differences in behaviour between strong and weak acids including the reaction with a reactive metal and difference in pH values by use of a pH meter, universal indicator or conductivity
33. understand that neutralisation reactions occur when H+(aq) and OH–(aq) form H2O(l)
34. understand that salts are formed in neutralisation reactions
35. sketch the pH titration curves of titrations using combinations of strong and weak acids with strong and weak alkalis
36. select suitable indicators for acid-alkali titrations, given appropriate data (pKa values will not be used)